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Avogadro's Number and the Mole
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Avogadro's Number and the Mole
Chemical Bridge Between Microscopic and Macroscopic Worlds
Sep 12, 2009 © Emily J. Foster
Chemists dealing with substance masses and tiny, microscopic entities use the mole - a
convenient unit that takes into account numbers of atoms and molecules.
Amedeo Avogadro (1776 - 1856) was an Italian physicist born in Turin to a noble family. He was noted for
his advances to molecular theory.
One of his most important contributions was to clearly distinguish
between atoms and molecules as separate entities, stating that gases are composed of molecules, while
such molecules are composed of atoms.
Avogadro's Law
Avogadro's Law states that equal volumes of all gases, under the same conditions of temperature and
pressure, contain equal numbers of molecules.
That is, the relationship between the masses of the same
volume of different gases is akin to the relationship between their respective molecular weights.
Gram-Atomic Weight - Avogadro's Number
Mass spectrometer measurements show that the atomic weight of, for example, magnesium is 24.31
atomic mass units (amu or u).
For the purpose of the present discussion, the u values can be approximated to the nearest whole number.
It can be shown experimentally that there is the same number of atoms in one gram-atom of any element.
That is, 1 g of hydrogen, 16 g of oxygen, 27 g of aluminium, 56 g of iron all contain the same number
of atoms - and that is called Avogadro's Number, a very large value of 6.02 x 10^23.
The original term gram-atom has since been replaced with the "mole", a concept explained next.
Avogadro's Number and the Mole
The new term mole is used to mean Avogadro's Number and is applied to a wide variety of entities
in the microscopic world, which are atoms and molecules.
So how does the mole relate to the macroscopic, "real" world? Well, a mole is defined as the number of particles in
exactly 12 grams of a particular isotope of carbon C-12. For any other element, a mole is the atomic
weight expressed in grams. For a compound, a mole is the formula weight (or molecular weight) in grams.
Examples of Usage of Mole
a) The atomic weight of oxygen is 16 u. Calculate the weight of one oxygen atom.
- The weight of 1 gram-atom of oxygen = 16 g. Since this is a mole of oxygen that contains 6.02 x 10^23 atoms, the weight
of a single atom of oxygen is 16/6.02 x 10^23 g, which equals 2.66 x 10^-23 g.
b) How many moles of aluminium are there in 90 g of aluminium?
- The weight of 1 gram-atom of aluminium = 27 g. Therefore, the moles of aluminium in 90 g is 90/27 = 3.3
c) Calculate the number of moles in one molecule of oxygen.
- The atomic weight of oxygen is 16 u and 1 u = 1.66 x 10^-24 g.
Weight of 1 molecule of oxygen = 2 x 16 x 1.66 x 10^-24 = 5.31 x 10^-23, therefore
Number of molecules in 32 g (i.e., in 1 mole) = 32/5.31 x 10^-23 = 6.02 x 10^23
The above are simple examples of the application of the mole concept forming a bridge from the very tiny,
microscopic size to the large, macroscopic size. For more further examples, the reader is directed to
the first quoted references (below).
References:
- Chemistry for Dummies. John T. Moore. Wiley Publishing, NJ. 2003
- Fundamental Chemistry. J.A. Saul. Hall's Book Store Pty Ltd, Melbourne, Australia. 1967.
The reader might be interested in another article
chemicals in the home , where everyone probably comes into contact with more chemicals and chemistry
in their own home than at any other place.
The copyright of the article Avogadro's Number and the Mole: Chemical Bridge Between Microscopic and Macroscopic
Worlds is owned by Emily J. Foster. Permission to republish in print or
online must be granted by the author in writing.
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